Osmium has a blue-gray tint and is the densest stable element, slightly denser than iridium.[5] Calculations of density from the X-ray diffraction data may produce the most reliable data for these elements, giving a value of 7004225620000000000♠22.562±0.009 g/cm3 for iridium versus 7004225870000000000♠22.587±0.009 g/cm3 for osmium.[6]

Osmium is a hard but brittle metal that remains lustrous even at high temperatures. It has a very low compressibility. Correspondingly, its bulk modulus is extremely high, reported between 7002395000000000000♠395 and 7011462000000000000♠462 GPa, which rivals that of diamond (7011443000000000000♠443 GPa). The hardness of osmium is moderately high at 7009400000000000000♠4 GPa.[7][8][9] Because of its hardness, brittleness, low vapor pressure (the lowest of the platinum-group metals), and very high melting point (the fourth highest of all elements), solid osmium is difficult to machine, form, or work.

The most common compound exhibiting the +8 oxidation state is osmium tetroxide. This toxic compound is formed when powdered osmium is exposed to air. It is a very volatile, water-soluble, pale yellow, crystalline solid with a strong smell. Osmium powder has the characteristic smell of osmium tetroxide.[20] Osmium tetroxide forms red osmates OsO
4(OH)2−
2 upon reaction with a base. With ammonia, it forms the nitrido-osmates OsO
3N−
.[21][22][23] Osmium tetroxide boils at 130 °C and is a powerful oxidizing agent. By contrast, osmium dioxide (OsO2) is black, non-volatile, and much less reactive and toxic.

Osmium pentafluoride (OsF5) is known, but osmium trifluoride (OsF3) has not yet been synthesized. The lower oxidation states are stabilized by the larger halogens, so that the trichloride, tribromide, triiodide, and even diiodide are known. The oxidation state +1 is known only for osmium iodide (OsI), whereas several carbonyl complexes of osmium, such as triosmium dodecacarbonyl (Os
3(CO)
12), represent oxidation state 0.[21][22][25][26]

In general, the lower oxidation states of osmium are stabilized by ligands that are good σ-donors (such as amines) and π-acceptors (heterocycles containing nitrogen). The higher oxidation states are stabilized by strong σ- and π-donors, such as O2−
and N3−
.[27]

Despite its broad range of compounds in numerous oxidation states, osmium in bulk form at ordinary temperatures and pressures resists attack by all acids and alkalis, including aqua regia.[28]

Osmium has seven naturally occurring isotopes, six of which are stable: 184
Os, 187
Os, 188
Os, 189
Os, 190
Os, and (most abundant) 192
Os. 186
Os undergoes alpha decay with such a long half-life ((2.0±1.1)×1015 years) that for practical purposes it can be considered stable. Alpha decay is predicted for all seven naturally occurring isotopes, but it has been observed only for 186
Os, presumably due to very long half-lives. It is predicted that 184
Os and 192
Os can undergo double beta decay but this radioactivity has not been observed yet.[29]

187
Os is the daughter of 187Re (half-life 7018143902655999999♠4.56×1010 years) and is used extensively in dating terrestrial as well as meteoricrocks (see rhenium-osmium dating). It has also been used to measure the intensity of continental weathering over geologic time and to fix minimum ages for stabilization of the mantle roots of continental cratons. This decay is a reason why rhenium-rich minerals are abnormally rich in 187
Os.[30] However, the most notable application of Os isotopes in geology has been in conjunction with the abundance of iridium, to characterise the layer of shocked quartz along the Cretaceous–Paleogene boundary that marks the extinction of the dinosaurs 66 million years ago.[31]

In 1803, Smithson Tennant analyzed the insoluble residue and concluded that it must contain a new metal. Vauquelin treated the powder alternately with alkali and acids[35] and obtained a volatile new oxide, which he believed to be of this new metal—which he named ptene, from the Greek word πτηνος (ptènos) for winged.[36][37] However, Tennant, who had the advantage of a much larger amount of residue, continued his research and identified two previously undiscovered elements in the black residue, iridium and osmium.[28][35] He obtained a yellow solution (probably of cis–[Os(OH)2O4]2−) by reactions with sodium hydroxide at red heat. After acidification he was able to distill the formed OsO4.[36] He named it osmium after Greekosme meaning "a smell", because of the ashy and smoky smell of the volatile osmium tetroxide.[38] Discovery of the new elements was documented in a letter to the Royal Society on June 21, 1804.[28][39]

Uranium and osmium were early successful catalysts in the Haber process, the nitrogen fixation reaction of nitrogen and hydrogen to produce ammonia, giving enough yield to make the process economically successful. At the time, a group at BASF led by Carl Bosch bought most of the world's supply of osmium to use as a catalyst. Shortly thereafter, in 1908, cheaper catalysts based on iron and iron oxides were introduced by the same group for the first pilot plants, removing the need for the expensive and rare osmium.[40]

Nowadays osmium is obtained primarily from the processing of platinum and nickel ores.[41]

Osmium is found in nature as an uncombined element or in natural alloys; especially the iridium–osmium alloys, osmiridium (osmium rich), and iridosmium (iridium rich).[35] In nickel and copper deposits, the platinum group metals occur as sulfides (i.e., (Pt,Pd)S)), tellurides (e.g., PtBiTe), antimonides (e.g., PdSb), and arsenides (e.g., PtAs2); in all these compounds platinum is exchanged by a small amount of iridium and osmium. As with all of the platinum group metals, osmium can be found naturally in alloys with nickel or copper.[43]

Osmium is obtained commercially as a by-product from nickel and copper mining and processing. During electrorefining of copper and nickel, noble metals such as silver, gold and the platinum group metals, together with non-metallic elements such as selenium and tellurium settle to the bottom of the cell as anode mud, which forms the starting material for their extraction.[46][47] In order to separate the metals, they must first be brought into solution. Several methods are available depending on the separation process and the composition of the mixture; two representative methods are fusion with sodium peroxide followed by dissolution in aqua regia, and dissolution in a mixture of chlorine with hydrochloric acid.[44][48] Osmium, ruthenium, rhodium and iridium can be separated from platinum, gold and base metals by their insolubility in aqua regia, leaving a solid residue. Rhodium can be separated from the residue by treatment with molten sodium bisulfate. The insoluble residue, containing Ru, Os and Ir, is treated with sodium oxide, in which Ir is insoluble, producing water-soluble Ru and Os salts. After oxidation to the volatile oxides, RuO
4 is separated from OsO
4 by precipitation of (NH4)3RuCl6 with ammonium chloride.

After it is dissolved, osmium is separated from the other platinum group metals by distillation or extraction with organic solvents of the volatile osmium tetroxide.[49] The first method is similar to the procedure used by Tennant and Wollaston. Both methods are suitable for industrial scale production. In either case, the product is reduced using hydrogen, yielding the metal as a powder or sponge that can be treated using powder metallurgy techniques.[50]

Neither the producers nor the United States Geological Survey published any production amounts for osmium. Estimations of the United States consumption date published from 1971,[51] which gives a consumption in the United States of 2000 troy ounces (62 kg), would suggest that the production is still less than 1 ton per year. In 2012, the estimated US production of osmium was 75 kg.[52]

Applications

Because of the volatility and extreme toxicity of its oxide, osmium is rarely used in its pure state, but is instead often alloyed with other metals. Those alloys are utilized in high-wear applications. Osmium alloys such as osmiridium are very hard and, along with other platinum-group metals, are used in the tips of fountain pens, instrument pivots, and electrical contacts, as they can resist wear from frequent operation. They were also used for the tips of phonographstyli during the late 78 rpm and early "LP" and "45" record era, circa 1945 to 1955. Although very durable compared to steel and chromium needle points, osmium-alloy tips wore out far more rapidly than competing but costlier sapphire and diamond tips and were discontinued.[53]

Osmium tetroxide has been used in fingerprint detection[54] and in staining fatty tissue for optical and electron microscopy. As a strong oxidant, it cross-links lipids mainly by reacting with unsaturated carbon–carbon bonds and thereby both fixes biological membranes in place in tissue samples and simultaneously stains them. Because osmium atoms are extremely electron-dense, osmium staining greatly enhances image contrast in transmission electron microscopy (TEM) studies of biological materials. Those carbon materials have otherwise very weak TEM contrast (see image).[24] Another osmium compound, osmium ferricyanide (OsFeCN), exhibits similar fixing and staining action.[55]

In 1898 an Austrian chemist Auer von Welsbach developed the Oslamp with a filament made of osmium, which he introduced commercially in 1902. After only a few years, osmium was replaced by the more stable metal tungsten. Tungsten has the highest melting point among all metals, and using it in light bulbs increases the luminous efficacy and life of incandescent lamps.[36]

The light bulb manufacturer Osram (founded in 1906, when three German companies, Auer-Gesellschaft, AEG and Siemens & Halske, combined their lamp production facilities) derived its name from the elements of osmium and Wolfram (the latter is German for tungsten).[59]

Like palladium, powdered osmium effectively absorbs hydrogen atoms. This could make osmium a potential candidate for a metal-hydride battery electrode. However, osmium is expensive and would react with potassium hydroxide, the most common battery electrolyte.[60]

Osmium has high reflectivity in the ultraviolet range of the electromagnetic spectrum; for example, at 600 Å osmium has a reflectivity twice that of gold.[61] This high reflectivity is desirable in space-based UV spectrometers, which have reduced mirror sizes due to space limitations. Osmium-coated mirrors were flown in several space missions aboard the Space Shuttle, but it soon became clear that the oxygen radicals in the low Earth orbit are abundant enough to significantly deteriorate the osmium layer.[62]

The only known clinical use of osmium appears to be for synovectomy in arthritic patients in Scandinavia.[63] It involves the local administration of osmium tetroxide (OsO4), which is a highly toxic compound. The lack of reports of long-term side effects suggest that osmium itself can be biocompatible, although this depends on the osmium compound administered. In 2011, osmium(VI)[64] and osmium(II)[65] compounds were reported to show anticancer activity in vivo, it indicated a promising future for using osmium compounds as anticancer drugs.[66]

Electron micrograph of (organic) plant tissue without (top) and with (bottom) OsO4 staining

Post-flight appearance of Os, Ag, and Au mirrors from the front (left images) and rear panels of the Space Shuttle. Blackening reveals oxidation due to irradiation by oxygen atoms.[67][68]

Precautions

Finely divided metallic osmium is pyrophoric[51] and reacts with oxygen at room temperature, forming volatile osmium tetroxide. Some osmium compounds are also converted to the tetroxide if oxygen is present.[51] This makes osmium tetroxide the main source of contact with the environment.

Osmium is usually sold as a minimum 99.9% pure powder. Like other precious metals, it is measured by troy weight and by grams. Its price in 2012 was about $400 per troy ounce (or about $13,000 per kilogram), depending on the quantity and its supplier.[71][72]

Tennant, S. (1804). "On Two Metals, Found in the Black Powder Remaining after the Solution of Platina". Philosophical Transactions of the Royal Society. 94: 411–418. doi:10.1098/rstl.1804.0018. JSTOR 107152.

MacDonell, Herbert L. (1960). "The Use of Hydrogen Fluoride in the Development of Latent Fingerprints Found on Glass Surfaces". The Journal of Criminal Law, Criminology, and Police Science. 51 (4): 465–470. doi:10.2307/1140672. JSTOR 1140672.

Chadwick, D. (2002). Role of the sarcoplasmic reticulum in smooth muscle. John Wiley and Sons. pp. 259–264. ISBN 0-470-84479-5.